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AP Chemistry Chapter 6 Outline
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AP Chemistry Chapter 6 Outline AP Chemistry Chapter 6 Outline Presentation Transcript

  • Chapter 6 Electronic Structure of Atoms Chemistry, The Central Science , 11th edition Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten John D. Bookstaver St. Charles Community College Cottleville, MO
  • Waves
    • To understand the electronic structure of atoms, one must understand the nature of electromagnetic radiation.
    • The distance between corresponding points on adjacent waves is the wavelength (  ) .
  • Waves
    • The number of waves passing a given point per unit of time is the frequency (  ) .
    • For waves traveling at the same velocity, the longer the wavelength, the smaller the frequency.
  • Electromagnetic Radiation
    • All electromagnetic radiation travels at the same velocity: the speed of light ( c ),
    • 3.00  10 8 m/s.
    • Therefore,
    • c = 
  • The Nature of Energy
    • The wave nature of light does not explain how an object can glow when its temperature increases.
    • Max Planck explained it by assuming that energy comes in packets called quanta .
  • The Nature of Energy
    • Einstein used this assumption to explain the photoelectric effect.
    • He concluded that energy is proportional to frequency:
    • E = h 
    • where h is Planck’s constant, 6.626  10 − 34 J-s.
  • The Nature of Energy
    • Therefore, if one knows the wavelength of light, one can calculate the energy in one photon, or packet, of that light:
    • c = 
    • E = h 
  • The Nature of Energy
    • Another mystery in the early 20th century involved the emission spectra observed from energy emitted by atoms and molecules.
  • The Nature of Energy
    • For atoms and molecules one does not observe a continuous spectrum, as one gets from a white light source.
    • Only a line spectrum of discrete wavelengths is observed.
  • The Nature of Energy
    • Niels Bohr adopted Planck’s assumption and explained these phenomena in this way:
      • Electrons in an atom can only occupy certain orbits (corresponding to certain energies).
  • The Nature of Energy
    • Niels Bohr adopted Planck’s assumption and explained these phenomena in this way:
      • Electrons in permitted orbits have specific, “allowed” energies; these energies will not be radiated from the atom.
  • The Nature of Energy
    • Niels Bohr adopted Planck’s assumption and explained these phenomena in this way:
      • Energy is only absorbed or emitted in such a way as to move an electron from one “allowed” energy state to another; the energy is defined by
      • E = h 
  • The Nature of Energy
    • The energy absorbed or emitted from the process of electron promotion or demotion can be calculated by the equation:
    where R H is the Rydberg constant, 2.18  10 − 18 J, and n i and n f are the initial and final energy levels of the electron.  E = − R H ( ) 1 n f 2 1 n i 2 -
  • The Wave Nature of Matter
    • Louis de Broglie posited that if light can have material properties, matter should exhibit wave properties.
    • He demonstrated that the relationship between mass and wavelength was
     = h mv
  • The Uncertainty Principle
    • Heisenberg showed that the more precisely the momentum of a particle is known, the less precisely is its position known:
    • In many cases, our uncertainty of the whereabouts of an electron is greater than the size of the atom itself!
    (  x ) (  mv )  h 4 
  • Quantum Mechanics
    • Erwin Schrödinger developed a mathematical treatment into which both the wave and particle nature of matter could be incorporated.
    • It is known as quantum mechanics .
  • Quantum Mechanics
    • The wave equation is designated with a lower case Greek psi (  ).
    • The square of the wave equation,  2 , gives a probability density map of where an electron has a certain statistical likelihood of being at any given instant in time.
  • Quantum Numbers
    • Solving the wave equation gives a set of wave functions, or orbitals , and their corresponding energies.
    • Each orbital describes a spatial distribution of electron density.
    • An orbital is described by a set of three quantum numbers .
  • Principal Quantum Number ( n )
    • The principal quantum number, n , describes the energy level on which the orbital resides.
    • The values of n are integers ≥ 1.
  • Angular Momentum Quantum Number ( l )
    • This quantum number defines the shape of the orbital.
    • Allowed values of l are integers ranging from 0 to n − 1.
    • We use letter designations to communicate the different values of l and, therefore, the shapes and types of orbitals.
  • Angular Momentum Quantum Number ( l ) f d p s Type of orbital 3 2 1 0 Value of l
  • Magnetic Quantum Number ( m l )
    • The magnetic quantum number describes the three-dimensional orientation of the orbital.
    • Allowed values of m l are integers ranging from - l to l :
    • − l ≤ m l ≤ l.
    • Therefore, on any given energy level, there can be up to 1 s orbital, 3 p orbitals, 5 d orbitals, 7 f orbitals, etc.
  • Magnetic Quantum Number ( m l )
    • Orbitals with the same value of n form a shell .
    • Different orbital types within a shell are subshells .
  • s Orbitals
    • The value of l for s orbitals is 0.
    • They are spherical in shape.
    • The radius of the sphere increases with the value of n.
  • s Orbitals
    • Observing a graph of probabilities of finding an electron versus distance from the nucleus, we see that s orbitals possess n − 1 nodes , or regions where there is 0 probability of finding an electron.
  • p Orbitals
    • The value of l for p orbitals is 1.
    • They have two lobes with a node between them.
  • d Orbitals
    • The value of l for a d orbital is 2.
    • Four of the five d orbitals have 4 lobes; the other resembles a p orbital with a doughnut around the center.
  • Energies of Orbitals
    • For a one-electron hydrogen atom, orbitals on the same energy level have the same energy.
    • That is, they are degenerate .
  • Energies of Orbitals
    • As the number of electrons increases, though, so does the repulsion between them.
    • Therefore, in many-electron atoms, orbitals on the same energy level are no longer degenerate.
  • Spin Quantum Number, m s
    • In the 1920s, it was discovered that two electrons in the same orbital do not have exactly the same energy.
    • The “spin” of an electron describes its magnetic field, which affects its energy.
  • Spin Quantum Number, m s
    • This led to a fourth quantum number, the spin quantum number, m s .
    • The spin quantum number has only 2 allowed values: +1/2 and − 1/2.
  • Pauli Exclusion Principle
    • No two electrons in the same atom can have exactly the same energy.
    • Therefore, no two electrons in the same atom can have identical sets of quantum numbers.
  • Electron Configurations
    • This shows the distribution of all electrons in an atom.
    • Each component consists of
      • A number denoting the energy level,
  • Electron Configurations
    • This shows the distribution of all electrons in an atom
    • Each component consists of
      • A number denoting the energy level,
      • A letter denoting the type of orbital,
  • Electron Configurations
    • This shows the distribution of all electrons in an atom.
    • Each component consists of
      • A number denoting the energy level,
      • A letter denoting the type of orbital,
      • A superscript denoting the number of electrons in those orbitals.
  • Orbital Diagrams
    • Each box in the diagram represents one orbital.
    • Half-arrows represent the electrons.
    • The direction of the arrow represents the relative spin of the electron.
  • Hund’s Rule
    • “ For degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized.”
  • Periodic Table
    • We fill orbitals in increasing order of energy.
    • Different blocks on the periodic table (shaded in different colors in this chart) correspond to different types of orbitals.
  • Some Anomalies
    • Some irregularities occur when there are enough electrons to half-fill s and d orbitals on a given row.
  • Some Anomalies
    • For instance, the electron configuration for copper is
    • [Ar] 4 s 1 3 d 5
    • rather than the expected
    • [Ar] 4 s 2 3 d 4 .
  • Some Anomalies
    • This occurs because the 4 s and 3 d orbitals are very close in energy.
    • These anomalies occur in f -block atoms, as well.